Topic 1: Structure of Water and Hydrogen Bonding
Quick recap: Chemistry foundations
Matter, element, and compound
Matter: anything that takes up space and has mass
Element: a substance that cannot be broken down into other substances by chemical reactions
Compound: a substance consisting of two or more different elements in a fixed ratio
Essential vs. trace elements in biology
Essential elements: among the 92 naturally occurring elements, ~20–25% are essential to survive and reproduce
CHOPN: Carbon (C), Hydrogen (H), Oxygen (O), Phosphorus (P), Nitrogen (N) – together ~96% of living matter
Trace elements: elements required in very small quantities
Task for students: look up essential elements and trace elements and answer questions in class
Recalling periodic trends lays groundwork for bonding concepts
Periodic table basics
Element symbol example: He for Helium
Atomic number: number of protons (top of the box in the diagram)
Atomic mass: sum of protons and neutrons, averaged over isotopes (bottom number in the box)
Groups vs. periods
Groups: vertical columns; elements in the same group have the same number of valence electrons
Periods: horizontal rows; elements in the same period have the same total number of electron shells
Think-pair-share prompt: review and take notes on what you remember about the periodic table for class discussion
Bonding fundamentals
Why atoms form bonds: to achieve stability by filling valence shells
Octet rule: atoms gain, lose, or share electrons to complete their valence shell
Valence shell: outermost electron shell involved in bonding
Lewis dot structures illustrate valence electrons and bonding
Example: carbon dioxide, CO₂, with two oxygen atoms bonded to carbon via double bonds
Definition of chemical bonds: attraction between atoms due to sharing or transferring valence electrons
Electronegativity: an atom’s ability to attract electrons
Trend: increases from left to right across a period; decreases down a group
Role in bond type determination (polarity of bonds)
Covalent vs ionic vs hydrogen bonds
Covalent bonds: sharing of electrons between two or more atoms (usually nonmetals)
Can form single, double, or triple bonds
Nonpolar covalent: equal sharing of electrons
Example: O₂, where O–O share electrons equally (same electronegativity)
Polar covalent: unequal sharing of electrons
Example: H₂O, oxygen is more electronegative than hydrogen, pulling electrons toward itself
Results in partial charges: O partially negative (δ−), H partially positive (δ+)
Ionic bonds: attraction between oppositely charged ions, usually metal + nonmetal
Electron transfer creates cations and anions
Examples: NaCl (sodium chloride)
Cation: positively charged ion; Anion: negatively charged ion
Hydrogen bonds: intermolecular attractions between molecules, not bonds within a single molecule
Occur when a partially positive hydrogen in one polar covalent molecule is attracted to a highly electronegative atom (commonly O or N) in another molecule
In water, hydrogen bonds form between the partially positive H of one H₂O and the partial negative O of a neighboring H₂O
Water properties: polarity, hydrogen bonding, and consequences
Water polarity
Polar covalent bonds within the water molecule (O–H bonds)
Oxygen is more electronegative than hydrogen, creating partial charges across the molecule
Hydrogen bonding in water
Intermolecular bonds between water molecules due to polarity
Visualized as pink dots in diagrams; facilitates many unique properties
Cohesion
Definition: attraction between molecules of the same kind (water–water)
Driven by hydrogen bonds; increases cohesion
Biological significance: supports transport of water and nutrients against gravity in plants; contributes to surface tension
Surface tension
Result of greater inward pull among surface water molecules
Creates a “film” at the air–water interface; helps support light objects like a paper clip on water
Adhesion
Attraction between water and other polar or charged surfaces
Enables water to cling to cell walls and to the xylem in plants, aiding upward movement
Capillary action
Upward movement of water driven by cohesion, adhesion, and surface tension
Occurs when adhesion to the container or plant tissue is stronger than cohesion between water molecules
Temperature control properties
High specific heat: water resists temperature changes due to hydrogen bonding
Energy is required to break hydrogen bonds; energy release occurs when bonds form
Real-world impact: moderates air and surface temperatures; stabilizes ocean temperatures; helps organisms resist internal temperature changes
High heat of vaporization (evaporative cooling)
Large amount of energy required for evaporation due to strong hydrogen bonding
Practical examples: sweating cools the body; leaves and climate stabilization in lakes/ponds
Density and phase behavior
Ice floats: solid water is less dense than liquid water
Hydrogen-bond network in ice forms a lattice, expanding as it freezes
Ecological significance: allows aquatic life to survive beneath frozen surfaces and supports life in polar regions
Water as a solvent
Solvent concept: water dissolves many substances due to polarity
“Like dissolves like”: polar solvents dissolve polar solutes and ionic compounds
Sugar in water example: sugar (polar) dissolves; water forms hydrogen bonds with sugars
Salt in water: Na⁺ and Cl⁻ become surrounded by water molecules via hydration shells; water’s partial charges interact with ions
Na⁺ is surrounded by partial negative oxygen atoms; Cl⁻ is surrounded by partial positive hydrogen atoms
Hydration and dissolution of ions
Dissolution equation (example): ext{NaCl}(s)
ightarrow ext{Na}^+(aq) + ext{Cl}^-(aq)Water’s dipole facilitates ion stabilization in solution
pH, acids, bases, and buffering
pH: measure of acidity/basicity of a solution
Acids: substances that release hydrogen ions (H⁺) in water
Bases: substances that accept H⁺ or release hydroxide ions (OH⁻) in water
Water autoionization (in simplified terms): ext{H}_2 ext{O}
ightleftharpoons ext{H}^+ + ext{OH}^-Buffer: a solution that resists pH changes when acids or bases are added
Biological relevance: buffers help maintain pH stability in organisms; water serves as a solvent for many buffers
Quick conceptual check (example from lecture)
Three-dimensional view of water can form up to four hydrogen bonds: each water molecule can form two bonds via its two hydrogen atoms and accept two bonds via its two lone pairs on oxygen
This means each H₂O can engage in up to four hydrogen bonds with neighboring water molecules
Applications and links to biology labs and coursework
Think-pair-share prompt (concept check): reflect on water properties, bonding, and their biological implications
Lab relevance: water properties stations; understanding pH and buffering is foundational for unit 2 and unit 4 in biology curricula
Real-world relevance
Plant physiology: cohesion, adhesion, and capillary action enable water transport in xylem and transpiration
Human physiology: evaporative cooling through sweating (high heat of vaporization)
Climate science: high specific heat of water moderates climate and stabilizes coastal ecosystems
Ethical/philosophical/practical implications
Water’s properties underpin life-supporting processes; anthropogenic changes to climate and water chemistry can disrupt these processes
Understanding solubility and buffering informs medicine, environmental policy, and industrial applications
Key formulas and equations (LaTeX)
Water structure and dissociation (conceptual):
ext{H}_2 ext{O}
ightleftharpoons ext{H}^+ + ext{OH}^-
Salt dissolution in water (example):
ext{NaCl}(s)
ightarrow ext{Na}^+(aq) + ext{Cl}^-(aq)
Carbon dioxide structure (Lewis dot/double bonds example):
CO₂ with two double bonds: ext{O}= ext{C}= ext{O}
Hydrogen bonding concept (between molecules): water–water interactions via partial charges created by polar O–H bonds
General solvent/solute terminology
Solvent: the dissolving agent (water in many biological contexts)
Solute: substance being dissolved (e.g., sugars, salts)
Sugar–water solution: example of a homogeneous mixture (solution)
Hydration shells: solvation of ions by water molecules (e.g., Na⁺ surrounded by O, Cl⁻ surrounded by H)
Connections to broader topics and study tips
Link to essential elements and macromolecules
Essential elements (C, H, O, P, N) contribute to the building blocks of macromolecules
Trace elements play critical, sometimes catalytic, roles in enzyme function and metabolism
Foundational principles applied across units
Polarity and bonding principles underpin biochemistry, cell biology, and physiology
Understanding pH, buffering, and solubility supports experiments in molecular biology and biochemistry
Practical implications for exams
Be able to explain differences between covalent (nonpolar vs polar), ionic, and hydrogen bonds
Describe how water’s properties (polarity, hydrogen bonding, cohesion/adherence, temperature regulation, density) arise from its molecular structure
Apply “like dissolves like” to predict solubility of solutes in water
Use simple chemical equations to illustrate dissolution and autoionization processes
Quick recap of key takeaways
Water’s polarity and hydrogen bonding drive its unusual properties: cohesion, adhesion, surface tension, high specific heat, high heat of vaporization, and ice’s lower density
Water acts as an excellent universal solvent for many biological molecules due to its polarity and ability to form hydration shells around ions
Bonding types (covalent, ionic, hydrogen) differ in how electrons are shared or transferred and in the scale (intramolecular vs intermolecular) at which they operate
pH and buffering are foundational for maintaining stable biological conditions; water participates in acid–base chemistry
Real-world implications span plant physiology, climate regulation, and human health, underscoring the importance of these concepts for biology and environmental science